09 Molecular Geometry and Polarity Online Quiz Questions

A central atom has one double bond, one triple bond, one single bond, and one lone pair. How many electron regions surround it, and what is the electron-domain geometry? Explain how you counted.

A central atom has three bonded atoms and one lone pair, represented as AX₃E. Determine its electron-domain geometry and molecular geometry. How should its bond angles compare with the ideal tetrahedral angle of $109.5^\circ$, and why?

For H₂O, account for the two O–H bonds and two lone pairs on oxygen. State the electron-domain geometry, the molecular geometry, and the approximate H–O–H bond angle. Explain why the angle is below the ideal tetrahedral value.

Rank lone pair–lone pair, lone pair–bonding pair, and bonding pair–bonding pair repulsions from strongest to weakest. Then explain how this hierarchy affects bond angles near a lone pair.

A central atom has one double bond and two single bonds, with no lone pairs. Use VSEPR to predict the basic arrangement around the central atom, then explain qualitatively why the three bond angles might not all equal the ideal trigonal-planar angle of $120^\circ$.

For AX₄E, identify the electron-domain geometry and molecular geometry. Where does the lone pair preferentially sit in the electron-domain arrangement, and what effect can it have on nearby bond angles?

An AX₃E₂ center has five electron regions. Derive its molecular geometry by identifying the preferred positions of the two lone pairs in a trigonal-bipyramidal arrangement. Describe how the resulting bonded-atom arrangement relates to the ideal angles.

For AX₂E₃, explain how the preferred placement of the three lone pairs in a trigonal-bipyramidal arrangement determines the molecular geometry and the ideal angle between the two bonded atoms.

A central atom is AX₅E. Determine its electron-domain geometry and molecular geometry, making clear how the lone pair affects the two names.

For AX₄E₂, determine the electron-domain and molecular geometries. Explain why the lone-pair positions lead to a planar arrangement of the four bonded atoms, and state the ideal angles between adjacent and opposite bonds.

Each C=O bond in CO₂ is polar. Explain how the molecule's geometry determines whether the two bond dipoles produce a net molecular dipole, and classify the molecule as polar or nonpolar.

CH₃Cl is tetrahedral. Explain why that shape does not make the molecule nonpolar, using the identities of the surrounding atoms and the vector sum of its bond dipoles to support your conclusion.

A central atom has three bonded atoms and two lone pairs. Use VSEPR theory to determine its electron-domain geometry and molecular geometry. Explain where the lone pairs preferentially go and give the typical bond angles.

A central atom is surrounded by five bonded atoms and one lone pair. Determine its electron-domain geometry and molecular geometry, and state the typical bond angles, noting whether they should be treated as exact.

A central atom has four bonded atoms and two lone pairs. Use VSEPR theory to identify the electron-domain geometry, the relative positions of the lone pairs, the molecular geometry, and its typical bond angles.

Compare NH₃ and H₂O using VSEPR theory. For each molecule, state its electron-domain geometry and molecular geometry, and explain why those two geometry names describe different arrangements.

NH₃ has a typical H–N–H angle of about 107°, while H₂O has a typical H–O–H angle of about 104.5°. Both have four electron regions around the central atom. Explain how the lone-pair counts and VSEPR repulsions help account for these angles compared with the ideal tetrahedral angle.

A central atom has one double bond to one atom, one single bond to a second atom, and one lone pair. Determine the number of electron regions, both geometries, and the expected bond-angle relationship to 120°. Explain how the double bond may affect the angle prediction.

Consider two molecules, each with two identical polar X–Y bonds. One has a linear XY₂ arrangement and the other a bent XY₂ arrangement. For each, explain whether the bond-dipole vectors cancel and predict whether the molecule is polar.

CO₂ and H₂O both contain polar bonds, yet one molecule is nonpolar and the other is polar. Explain this difference by identifying each shape and describing the vector sum of its bond dipoles.

Explain why BCl₃ is nonpolar while CH₃Cl is polar. In your comparison, identify each molecule’s geometry and explain how the identities and arrangement of the surrounding atoms affect the bond-dipole vector sum.

A molecule is known to be trigonal planar and to contain polar bonds, but the identities of its three surrounding atoms are not specified. Can you determine whether the molecule is polar? Explain what information is missing and how it affects the dipole-vector sum.

Describe a reliable step-by-step method for deciding whether a molecule is polar. Explain why stopping after identifying polar bonds would not be sufficient.

A central atom has two identical bonded atoms and three lone pairs, giving an AX₂E₃ arrangement. Determine the electron-domain geometry, preferred lone-pair positions, molecular geometry, and typical bond angle. If the two bonds are polar, also determine whether their dipoles cancel and justify your conclusion.

A central atom has one double bond to one atom, one single bond to a second atom, and one lone pair. Using VSEPR, identify its electron-domain geometry and molecular geometry, and state how its bond angle compares with $120^\circ$. Explain your reasoning.

For a central atom with five electron regions, two bonded atoms, and three lone pairs ($AX_2E_3$), identify the electron-domain geometry and molecular geometry, and give the bond angle. Explain how the region count leads to your result.

Compare $\mathrm{NH_3}$ and $\mathrm{H_2O}$. For each molecule, state its electron-domain geometry, molecular geometry, and approximate bond angle. Explain why their angles differ even though each central atom has four electron regions.

Using bond-dipole vectors, explain why $\mathrm{CO_2}$ is nonpolar while $\mathrm{H_2O}$ is polar. Include the role of each molecule’s geometry in your explanation.

Compare $\mathrm{BCl_3}$ and $\mathrm{CH_3Cl}$. For each, state its geometry and whether it is polar or nonpolar, then explain why geometry by itself is not enough to predict the outcome in both cases.

Imagine a molecule with three bond-dipole vectors of equal magnitude arranged symmetrically in a trigonal plane. Predict whether their sum is zero. Then predict what happens if one vector becomes larger while all three directions stay fixed. Explain both conclusions using vector addition.